Acids, Bases, and pH

Definitions and Key Concepts

Acids and bases are fundamental concepts in chemistry and biology, including microbiology, as they influence enzyme activity, cellular processes, and buffer systems in biological fluids.

• Acid: A substance that donates protons (H+) in aqueous solution.

• Base: A substance that accepts protons (H+).

• Conjugate Acid-Base Pair: Each acid is linked to a conjugate base on the other side of the equation.

• pH: A measure of the hydrogen ion concentration in a solution, indicating its acidity or alkalinity.

Key Equation:

Where [H+] is the concentration of hydrogen ions in the solution.

Calculating pH of Strong Acids and Bases

Strong Acids

Strong acids completely dissociate in water. The concentration of hydrogen ions in a monoprotic strong acid will be the same as the concentration of the acid.

• For HCl and HNO3, [H+] = concentration of acid.

• Example: For 0.1 mol dm-3 HCl, pH = -log(0.1) = 1.00

Finding [H+] from pH:

Example: If pH = 1.35, mol dm-3

Strong Bases

Strong bases also completely dissociate into their ions. For NaOH, [OH-] = concentration of base.

• To work out pH, first calculate [OH-], then use the ionic product of water to find [H+].

Ionic Product of Water:

At 25°C, mol2 dm-6

Weak Acids and Bases

Weak Acids

Weak acids only partially dissociate in water, establishing an equilibrium mixture.

• Acid dissociation constant (Ka): Measures the strength of a weak acid.

The larger the Ka, the stronger the acid.

pKa:

Calculating pH of a Weak Acid

Assumptions for calculation:

• The acid is weak, so only a small amount dissociates.

• The initial concentration of acid is approximately equal to the equilibrium concentration.

Example: For ethanoic acid, mol dm-3

Buffer Solutions

Definition and Function

A buffer solution resists changes in pH when small amounts of acid or alkali are added. Buffers are crucial in biological systems, such as blood, to maintain a stable pH environment.

• Acidic buffer: Contains a weak acid and its salt (conjugate base).

• Basic buffer: Contains a weak base and its salt (conjugate acid).

How Buffer Solutions Work

• If acid is added, the conjugate base in the buffer neutralizes added H+ ions.

• If alkali is added, the weak acid in the buffer neutralizes added OH- ions.

Buffer Equation:

Or, rearranged for pH:

Buffering Action in Blood

The carbonic acid–bicarbonate equilibrium acts as a buffer in the control of blood pH:

This buffer maintains blood pH between 7.35 and 7.45.

Titration Curves and Indicators

Titration Curves

Titration curves show how pH changes as acid or base is added. The shape of the curve depends on the strength of the acid and base involved.

• Strong acid + strong base: Steep vertical region at equivalence point (pH ~7).

• Weak acid + strong base: Buffer region before equivalence point; equivalence point at pH > 7.

Choosing an Indicator

• Indicators are weak acids or bases that change color at a specific pH range.

• The indicator should have a color change range that matches the steep part of the titration curve (equivalence point).

Indicator Equation:

Enthalpy Change of Neutralisation

The standard enthalpy change of neutralisation is the enthalpy change when solutions of an acid and an alkali react together under standard conditions to produce 1 mole of water.

• For strong acids and bases, the value is typically between -56 and -58 kJ mol-1.

Tables

Comparison of Acid and Base Strengths

Summary of Key Equations

Additional info:

• These concepts are foundational for understanding microbial growth environments, enzyme activity, and metabolic processes in microbiology.

• Buffer systems are especially important in maintaining physiological pH in biological systems.